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Chapter 3 - Atoms:
The Building Blocks of Matter
• There were two schools of thought of
the composition of the cosmos…
– is everything in the universe
continuous and infinitely divisible
– Or, is there a limit to how small you
can get?
• Particle theory was not the most
popular early opinion, but was
supported as early as Democritus
in ancient Greece.
From Philosophy to Science
• Democritus proposed that all the
matter is composed of tiny particles
called “Atomos”
– These “particles” were thought to be
indivisible
• Aristotle did not accept Democritus’
atom, he was of the “matter is
continuous” philosophy
– Because of Aristotle’s popularity
his theory was adopted as the
standard
From Philosophy to Science
• By the 1700’s nearly all chemists had
accepted the modern definition of an
element as a particle that is indivisible
• It was also understood at that time that
elements combine to form compounds
that are different in their properties than
the elements that composed them
– However, these understandings
were based on observations
not empirical evidence
From Philosophy to Science
• There was controversy as to whether
elements always combine in the same
proportion when forming a particular
compound.
– In the 1790’s, chemistry was
revolutionized by a new emphasis
on quantitative analysis because
of new and improved balances
• This new technology led
to the discovery of some new
scientific understandings
From Philosophy to Science
• The Law of Conservation of Mass:
– States that mass is neither created nor
destroyed during ordinary chemical
reactions or physical changes.
– Which means the total mass of the
reactants must equal the total mass
of the products.
From Philosophy to Science

+
Carbon, C
Mass x
Oxygen, O
Mass y
Carbon Monoxide, CO
Mass x + Mass y

Carbon Monoxide, CO
Mass x + Mass y
+
Carbon, C
Mass x
Oxygen, O
Mass y
• The Law of Definite Proportions:
– The fact that a chemical compound
contains the same elements in exactly
the same proportions by mass
regardless of the size of the sample
or the source of the compound
• NaCl is NaCl no matter if it is table
salt (small crystals) or rock salt
(large crystals)
From Philosophy to Science
• The Law of Multiple Proportions:
– If 2 or more different compounds
are composed of the same 2 elements,
then the ratio of the masses of the
2nd element combined with a certain
mass of the 1st element is always
a ratio of small whole numbers
From Philosophy to Science
+
Carbon
1
=
Oxygen
1
+
Carbon Monoxide,
1:1
=
Carbon
Oxygen
1
2
Carbon Dioxide,
1:2
• In 1808, John Dalton proposed an
explanation for each of the proposed
laws
– He reasoned that elements were
composed of atoms & that only whole
#’s of atoms can combine to form
compounds
– His ideas are now called the Atomic
Theory of Matter
Atomic Theory
1. All matter is composed of
extremely small particles called
atoms.
2. Atoms of a given element are
identical in size, mass, and
other properties; atoms of
different elements differ in size,
mass, & other properties.
ELEMENT
2
ELEMENT
3
ELEMENT
4
Atomic Theory
3. Atoms cannot be
subdivided, created, or
destroyed
4. atoms of different
elements combine in
simple whole # ratios to
form chem compds
5. in chemical rxns, atoms
are combined,
separated, or rearranged
Atomic Theory
+
+
• Through these statements, evidence
could be gathered to confirm or
discount its claims
– Not all of Dalton’s claims held up
to the scrutiny of experimentation
– Atoms CAN be divided into even
smaller particles
– Not every atom of an element
has an identical mass
Atomic Theory
• Dalton’s Atomic Theory of Matter has
been modified.
• What remains unchanged is…
1.All matter is composed of atoms
2.Atoms of any one element differ in
properties from atoms of another
element
Atomic Theory
One of the disputed statements of Dalton
was that atoms are indivisible.
– In the 1800’s it was determined that
atoms are actually composed of
several basic types of smaller particles
– it’s the number and arrangement of
these particles that determine the
atom’s chemical properties.
Atomic Theory
• The definition of an atom that emerged
was: the smallest particle of an
element that retains the chemical
properties of that original
element.
• All atoms consist of 2 regions that
contain the subatomic particles
– The nucleus
– The electron cloud around the
nucleus
Atomic Theory
• The nucleus is a very small region
located near the center of the atom
– In every atom the nucleus contains
at least 1 proton, which is positively
charged particle and usually
contains 1 or more neutral
particles called neutrons
Atomic Structure
• The electron cloud is the region that
surrounds the nucleus
– This region contains 1 or more
electrons, which are negatively charged
subatomic particles
– The volume of the
electron cloud is much
larger than the nucleus
Atomic Structure
• With the exception of Hydrogen, every
nucleus contains 2 kinds of particles
protons and neutrons
– they make up the mass of the atom (Mass
Number = Protons + Neutrons)
• Proton has a charge equal to but
opposite of the charge of an electron.
– Atoms are neutral because they contain
equal #’s of protons & electrons
Atomic Structure
• The atoms of different elements differ
in the # of protons in their nuclei and
therefore in their positive charge
– The # of protons the atom contains
determines the atom’s identity, also
known as atomic number.
• Only Oxygen contains 8 protons
• Only Fluorine contains 9 protons
• Only Neon contains 10 protons
Structure of the Atom
Particle Symbol Charge
Mass
Number
Relative Mass
(amu)
Actual Mass
(kg)
Electron
e-
-1
0
0.0005486
9.109 x 10-31
Proton
Neutron
p+
n0
+1
0
1
1
1.007276
1.008665
1.673 x 10-27
1.675 x 10-27
Ch 3.3: Atomic Number
• Elements are identified by the number of
PROTONS they contain.
• The “atomic number” of an element is
the number of protons in the nucleus
– PROTONS IDENTIFIES AN ELEMENT!!!
• # protons in an atom = # electrons
–Why? Because atoms are neutral!
Complete Symbol
Superscript →
Mass
number
Subscript →
Atomic
number
X
# OF PROTONS
+
# OF NEUTRONS
MASS
NUMBER
35
ATOMIC
NUMBER
17
Cl
NUMBER OF
PROTONS
Mass Number
Mass number is the number of protons
and neutrons in the nucleus of an
+ + n0
Mass
#
=
p
isotope:
p+
n0
e-
Mass #
8
10
8
18
Arsenic - 75
33
42
33
75
Phosphorus - 31
15
16
15
31
Element
Oxygen -
Mass number
18
Practice Problems
(1)Find the # of e-, p+ and n0 for sodium.
(mass # = 23)
Atomic # = 11 = # e- = # p+
# neutrons = 23-11 = 12
2) Find the # of e-, p+ and n0 for uranium.
(mass # = 238)
Atomic # = 92 = # e- = # p+
# neutrons = 238-92 = 146
Check for understanding:

If an element has 91 protons
and 140 neutrons find the:
a) Atomic number 91
b) Mass number 231
c) number of electrons 91
d) element name protactinium
Isotopes
• An isotope refers to atoms that have
the same # of protons, but a different
number of neutrons.
• Because of this, they have different
mass #’s.
Ex---> (1) Carbon-12 & Carbon-13
(2) Chlorine-35 & Chlorine-37
(Isotopes: The # after the name is the
mass #.)
EXAMPLE OF AN ISOTOPE
ATOMIC MASS
Cl
Cl
35
37
17
17
18
NEUTRONS
20 NEUTRONS
ATOMIC NUMBER
Question #1

Find each of these:
a) Atomic number
b) Mass Number
c) number of protons
d) number of neutrons
e) number of electrons
80
35
Br
Question #2

If an element has an atomic
number of 34 and a mass
number of 78, what is the:
a) number of protons
b) number of neutrons
c) number of electrons
d) complete symbol
Atomic Mass
12
• Units = atomic mass unit (amu)
• The atomic masses listed in the
Periodic Table are a “weighted
average” of all the isotopes of the
element.
Weighted Average
Practice Problems:
(1) In chemistry, chlorine has 2 isotopes:
Cl-35 (75.8% abundance) Cl-37 (24.23 % abundance)
What is the weighted average atomic mass of chlorine?
35 x 0.758 = 26.53 amu
37 x 0.2423 = + 8.965 amu
35.495 amu
This rounds to 35.5 amu
Add them up!!!
Relating Mass Numbers to Atoms
• The Mole: the amount of a substance that
contains as many particles as there are
atoms in exactly 12 grams of carbon-12.
• Avogadro’s Number: the number of
particles in exactly one mole of a pure
substance = 6.022 x 1023.
• Molar Mass: the mass of one mole of a
pure substance. Units = g/mol
• This is when we get to use dimensional
analysis!
• The conversion factors we need are:
1mol
23
6.022 x10 atoms
6.022 x10 atoms
1mol
and of course…molar mass
___ g
1mol
1mol
___ g
23
Gram to Mole Conversions
___ g
1mol
Mass of
Element in
Grams
23
6.022 x10 atoms
1mol
Number of
Moles of
Element
1mol
___ g
Number of
Atoms of
Element
1mol
6.022 x1023 atoms
Practice Problem
• ALWAYS USE PARANTHESES AROUND
YOUR CONVERSION FACTORS!!
• You have 3.50 mol of Copper. What is it mass in
grams?
• You have 28.55 grams of Carbon. How many
atoms is this?
• You have 4.85 x 1030 atoms of Magnesium. How
many moles is this?